Exceptions to the octet rule
** For a time it was believed that all compounds obeyed the Octet rule or the Rule of eight. However, it gradually became apparent that quite a few molecules had non-octet structures. Atoms in these molecules could have number of electrons in the valence shell short of the octet or in excess of the octet.
** Some important examples are:
(1) Four or six electrons around the central atom
** A stable molecule as of beryllium chloride, BeCl2, contains an atom with four electrons in its outer shell.
** The compound boron trifluoride, BF3, has the Lewis structure :
The boron atom has only six electrons in its outer shell.
Beryllium chloride and boron trifluoride are referred to as electron-deficient compounds
(2) Seven electrons around the central atom
** There are a number of relatively stable compounds in which the central atom has seven electrons in the valence shell. A simple example is chlorine dioxide, ClO2.
The chlorine atom in ClO2 has seven electrons in its outer shell.
** Methyl radical (CH3) has an odd electron and is very short lived. When two methyl free radicals collide, they form an ethane molecule (C2H6) to satisfy the octet of each carbon atom.
** Any species with an unpaired electron is called a free radical.
(3) Ten or more electrons around the central atom
** Non-metallic elements of the third and higher periods can react with electronegative elements to form structures in which the central atom has 10, 12 or even more electrons. The typical examples are PCl5 and SF6.
** The molecules with more than an octet of electrons are called superoctet structures.
** In elements C, N, O and F the octet rule is strictly obeyed because only four orbitals are available (one 2s and three 2p) for bonding. In the elements P and S, however, 3s, 3p, and 3d orbitals of their atoms may be involved in the covalent bonds they form. Whenever an atom in a molecule has more than eight electrons in its valence shell, it is said to have an expanded octet.
Variable Valence
** Some elements can display two or more valences in their compounds.
** The transition metals belong to this class of elements. The Electronic Structure of some of these metals is given below:
** Most of the transition metals have one or two outer-shell electrons and they form monovalent or bivalent positive ions. But because some of the d electrons are close in energy to the outermost electrons, these can also participate in chemical bond formation. Thus transition metals can form ions with variable valence.
** For example, copper can form Cu1+ and Cu2+ ions and iron can form Fe2+and Fe3+ ions.
** The complete electronic configuration of an iron atom is:
It can form Fe2+by losing two 4s electrons,
When iron loses two 4s electrons and one of the three 3d electrons, if forms Fe3+ion
** Copper form Cu1+ and Cu2+ ions by losing one 4s electron, and one 4s and 3d electron respectively
** It may be noted that the structures of Fe2+, Fe3+, Cu1+, Cu2+, Cr3+, etc., are not isoelectronic with any of the noble gases, and hence the d electrons being unstable are available for bond formation.
** The atoms and ions that have the same number of electrons are said to be Isoelectronic.
Reference: Essentials of Physical Chemistry /Arun Bahl, B.S Bahl and G.D. Tuli / multicolour edition.
